SCH3U Exam Review1 11U CHEMISTRY EXAM REVIEW QUESTIONS (Part 1) UNITS 1 -3 1. Classify the following as chemical or physical properties of matter. a) silver does tarnish b) gold is a conductor of electricity c) potassium iodide dissolves in water d) Oxygen supports combustion e) Mercury metal can evaporate f) Lead has a high density 2. Identify each of the changes below as chemical (C) or physical (P). a) twisting copper wire into a coil c) melting wax b) burning coal d) fermentation of grapes to make wine 3. Classify each of the following as a pure substance (PS), a solution (S), or a mechanical mixture (MM). a) plain Jell-0 d) Raisin Bran cereal b) neon gas e) calcium carbonate c) vinegar f) air 4. a) What is the name for the smallest particle of a covalently-bonded compound ? _________________ b) Only pure substances are homogeneous, true or false? c) Explain your answer to b). 5. a) J.J Thomson discovered the __________________ while studying ___________________ rays b) Very briefly describe Thomson's model of the atom. 6. Rutherford did an experiment that disproved Thomson's model of the atom. In his experiment Rutherford shot a beam of ______________ particles at a thin sheet of _________. Most the these particles ___________________ and this proved that 99.99% of the atom is consists of _________ ______________. However, about 1 in 20,000 of the particles were_____ ____________ and this proved the existence of a very dense ______________________ in the atom. Based on the fact that the positive charge on an atom is a multiple of that of hydrogen, Rutherford concluded that this structure contained a subatomic particle called the ____________________. Later experiments by ________________________ discovered the last major subatomic particle found in this region, the ____________________. 7. Consider the following atoms : 16 8X, 17 9Y, 17 8Z Which of these atoms (if any) are isotopes of the same element? Explain your reasoning. 8. Draw Bohr-Rutherford diagrams for the following isotopes of each element. a) Sodium-23 (2311Na) b) Phosphorus-31 (3115P) 9. Use the information given below to calculate the average atomic mass of element X. ISOTOPE 302 X 304 X 306 X ABUNDANCE (%) 12.64 18.23 69.13 MASS (u or g/mol) 302.04 304.12 305.03 SCH3U Exam Review 2 10. . alkali metals. Cl or Br f) Circle which has the greatest metallic character and state why. a) Circle which is bigger and state why . Element X is composed of isotopes X-56 (mass 56. polar covalent and ionic bonds. Consider the following electron configurations for neutral atoms: i) 2. 1 a) Which of these would you expect to be an alkaline earth metal? ______________ b) List the five atoms in order of increasing atomic radii. 8. _______________________ c) Which of these would you expect to have the lowest ionization energy ? _______________ d) Which of these would you expect to be in the same family? _______________ 15. S or Cl e) Circle which has the greater electron affinity and state why.0 u) and X-59 (mass 59. noble gases. 7 ii) 2. Which of the following ions would you expect to.0 u). Explain the difference between covalent. Li or Na b) Circle which is bigger and state why .8. 8. All of the trends on the periodic table can be explained by two factors : i) the number of occupied Bohr energy levels (orbits) ii) the number of protons in the nucleus (nuclear charge) Answer each of the following using these 2 factors. 12. O or F c) Circle which is bigger and state why . halogens.3 u? 11. Be or Mg 16. be stable? Mg2+ K- B3- S2- 14. 8. transition metals. How many valence electrons are there in: a) calcium b) fluorine 13. What is the % abundance of each of these isotopes if the atomic mass is 57. Na+ or Ne d) Circle which has the higher ionization energy and state why. 8 iv) 2. On the periodic table label the following groups of elements: metals. nonmetals. 8. 2 v) 2. metalloids and the inner transition metals. 8 iii) 2. 8. Write the ionization equations for each of the following: a) calcium b) phosphorus 17. Draw a Lewis structure and name the shape for these compounds.SCH3U Exam Review 3 18. a) BrCl b) NI3 c) Al2O3 d) Cs2S 19. Compound Name carbon tetrachloride Formula Lewis Structure Name of 3D Shape nitrogen nitrogen trifluoride hydrogen peroxide boron trihydride 20. Determine the bond type in the following compounds using electronegativity values. Complete the following table. Formula PbS Name nitric acid SnSO3•6H2O plumbic iodide SF6 neon Na2O2 Al(IO4)3 iron (III) hydrogen sulfite chlorine NaBrO potassium bromite Ca(HCO3)2 KH2PO3 copper (II) sulfate nonahydrate manganese (II) hypochlorite hydrosulfuric acid HClO2 (aq) BaO2 KNO2 manganese dioxide Fe(OH)2 (ous/ic) sulfuric acid NH4OH silver sulfide Mg3P2 lead (IV) perchlorate Au(OH)3 magnesium phosphite MgO Name Formula . The following observations were made when the signle displacement reactions were attempted: a) Z (s) + WA (aq) b) W (s) + XA (aq) c) X (s) + ZA (aq) d) W (s) + YA (aq) → → → → ZA (aq) + W (s) NR XA (aq) + Z (s) WA (aq) + Y (s) . b) Write the type of reaction (synthesis. decomposition. a) b) c) d) e) Ca3(PO4)2 + SiO2 Al2C6 + H2O → FeCl2 + KMnO4 BiO2 + H2 FeS + O2 + C → P4 + CaSiO4 + CO Al(OH)3 + C2H2 + HCl → FeCl3 + KCl + MnCl2 + H2O → Bi + H2O → Fe2O3 + SO2 22. In an experiment designed to compare the reactivity of 4 metals (W. predict if a reaction will occur. If no reaction occurs. X. single displacement. neutralization) beside each. a) water → b) silver + copper (II) sulfate → c) carbon dioxide + water → d) barium chloride + sodium sulfate → e) butane (C4H10) + oxygen → f) sodium hydroxide + nitric acid → 24.SCH3U Exam Review 4 21. combustion. a) Complete these reactions and write a balanced chemical equation. complete the word equation. a) b) c) tin (s) + potassium iodide (aq) fluorine (aq) + sodium chloride (aq) sodium (s) + calcium sulfate (aq) 23. the metallic elements were reacted with aqueous solutions containing their ions and an anion A. For each pair of reactants. write NR. double displacement. Y and Z). If a reaction occurs. Balance the following reactions. given the data in problem b)? d) What mass of water is produced when 105 g of hydrogen reacts with excess oxygen? 6. Calculate the molecular formula of hydroquinone. a) limiting and excess reactant b) actual yield and theoretical yield c) empirical and molecular formula 2. Calculate the percent composition of sodium thiosulfate. 4. A crystal is added to a solution. What mass of sodium carbonate must be used to produce 10. Explain the difference between each pair of terms.7% oxygen.0 g of oxygen gas and 300 g of propane are mixed and allowed to react as shown below. Determine the empirical formula for the compound.36 L of carbon dioxide at 24°C and 103 kPa according to the following neutralization reaction? HCl (aq) + Na2CO3 (s) → NaCl(aq) + CO2 (g) + H2O (l) 7. One of the reactions used in the smelting of copper ores to produce copper and sulfur dioxide involves reacting copper(I) oxide with copper(I)sulfide. determine the volume of water vapour formed at 116 kPa and 120°C. d) Determine the percentage yield of copper. It has a molecular mass of 110 g/mol and a composition of C3H3O. how many grams of oxygen must be consumed? c) How many moles of hydrogen must be used.00 moles of water are produced. 5. Describe what happens to the crystal if the solution is unsaturated. saturated or supersaturated. .6% potassium. Hydroquinone is an organic compound commonly used as a photographic developer. Use this equation for all three problems: 2 H2 (g) + O2 (g) → 2H2O (g) a) How many moles of water are produced when 5. b) Determine the limiting reagent. 9. A compound contains 56. When 250 kg of copper (I) oxide is heated with 129 kg of copper (II) sulfide.00 moles of oxygen are consumed? b) If 3. Cu2O (s) + Cu2S (s) → Cu (s) + SO2 (g) a) Write the balanced chemical equation for the reaction. 8.7% carbon and 34. If 129. C3H8 (g) + O2 (g) → CO2 (g) + H2O (g) 8. Na2S2O3. 3.SCH3U Exam Review 5 11U CHEMISTRY EXAM REVIEW (Part 2) UNITS 4 – 6 1. 285 kg of copper is recovered. c) Calculate the theoretical yield using stoichiometry. 50% by mass acetic acid solution if the density is 1. Complete balanced chemical equations for the following precipitation reactions.02 g/mL.8 ppm.75 mol/L solution? 13. 25 g of sodium chloride is dissolved in 100 g of water.15 g/mL) 12.) 15. Complete a balanced chemical equation. total ionic equation and net ionic equation. 16. What is the concentration in: a) % by mass b) mol / L (the density of the solution is 1.000 L pool? (DH2O 1. Acetic acid has the formula CH3COOH.00 g/mL = 1. What mass of lead would there be in a 60. What volume of concentrated phosphoric acid is needed to make 8.50 mol/L solution? 14. a) Fe(NO3)2 (aq) + NaOH (aq) → b) AgNO3 (aq) + MgCl2 (aq) → 18. a) H3PO4 (aq) + NaOH (aq) → b) HI (aq) + Na2CO3 (aq) → 20. Consider the solubility curves.00 kg/L.0 mol/L. Pb(NO3)2 (aq) + KI (aq) → 19. Include the sate symbols s or aq. Writa a balanced dissociate equation for these salts. . Describe the Arrhenius and Bronsted-Lowry definitions of an acid and a base.SCH3U Exam Review 6 10. a) KCl(s) b) Ba(NO3)2 (s) c) (NH4)2SO3 (s) 17.00 L of a 1. What mass of lead (II) nitrate is required to make 4.5 L of a 0. Complete balanced neutralization reactions. The water in a swimming pool has tested positive for lead at a concentration of 4. a) What is the solubility of KNO3 at 50oC? b) What must the temperature be to create a saturated solution of KNO3 using 50 g of the salt and 50 g of water? c) What mass of KNO3 will precipitate if a saturated solution of KNO3 at 60oC is cooled to 10oC? 11. Concentrated phosphoric acid has a concentration of 18. Calculate the concentration (in mol/L) of a 7. 2 L of carbon dioxide gas at 25oC and 97. determine the molar mass. A 50.SCH3U Exam Review 7 21.5 mL of sodium hydroxide. What is the hydronium concentration of a hydrobromic acid solution with a pH of 0.520 mol/L sodium hydroxide are needed to neutralize 100 mL of a 7. 22. 29. Calculate the concentration of the base. bp = 80. 27.5 L at a pressure of 770 mmHg. 33. Calculate the temperature required to change the volume to 42. A student collects 300 mL of hydrogen over water at 105. A sample of gas has a volume of 54. assuming pressure and temperature remain constant? 36.2 L of carbon dioxide gas at STP c) moles in 45.05 mol/L HCl b) 0. A student collected 245 mL of an unknown gas X over water at an atmospheric pressure of 108. If the mass of the gas is 0. 23.0 kPa. 28. What must the pressure be changed to in order to make the volume 0.7? 26.2 g of carbon dioxide gas at STP f) 45.50 mol/L hydrobromic acid solution? Include a balanced chemical equation.6 kPa and a temperature of 20oC.700 L? 31. What is the density of methane gas (CH4) at 20oC and 115 kPa? 35. . 34.g. 32.0 L at 22oC.g. Write an equation that demonstrates how pure water can act as both a Bronsted Lowry acid and Bronsted Lwory base at the same time. e) 45.1oC). acetic acid) 24. Sketch and label a heating curve for benzene (mp = 5. Distinguish between strong acids (e. how many litres of nitrogen trifluoride gas will be produced.0 mL of 0. If 1.2 L of carbon dioxide gas at STP b) molecules in 45.00005 mol/L HNO3 25. Include a balanced chemical equation. How many moles of the gas were collected? Be sure to correct for vapour pressure.0 L volume of gas is at 42oC and 87 kPa. List 3 properties of acids and 3 properties of bases.2 L of oxygen gas at STP. 10. The pressure on 925 mL of a gas is 120 kPa.5oC.150 mol/L sulfuric acid is titrated with 24. d) moles in 45. Calculate the number of: a) moles in 45.7 kPa and 25oC. What will the volume be if it is cooled to -50oC? 30.00 L of nitrogen gas is reacted with 3.0 L of fluorine gas. hydrochloric acid) and weak acids (e. What volume of 0. What is the pH of the following solutions? a) 0.7802 g.2 g of oxygen gas at STP. 18. CCl4. H2O2. N2. Mg2+ and S2. P. I d) 1. 2. Chadwick. deflected. Molecule 5.0.are stable ions (they have afull valence shell) 14. PS. C 19. a) iv b)smallest ii < iii< I < iv < v largest 15. 5 d) ii and iii are both noble gases a) Na (more energy levels) b) O (less nuclear charge) c) Ne (same electronic arrangement. nucleus. S. P 2. alpha (helium nuclei). C. P. proton. PS. tetrahedral.5 C/PC c) 2. P. trigonal pyramidal. nucleus. gold. linear. Polar colvalent: unequal sharing of a pir of electrons (have bond dipoles). neutron 7.8. C. 304. electron. a) Ca → Ca2+ + 2e- b) P + 3e. angular. Formula PbS SnSO3•6H2O SF6 Na2O2 Al(IO4)3 NaBrO Ca(HCO3)2 MgO KH2PO3 Name Lead (II) sulfide Tin (II) sulfite hexahydrate Sulfur hexafluoride Sodium peroxide Aluminum periodate Sodium hypobromite Calcium hydrogen carbonate Magnesium oxide HClO2 (aq) KNO2 Fe(OH)2 (ous/ic) NH4OH Potassium dihdrogen phosphate Chlorous acid Potassium nitrite ferrous hydroxide Ammonium hydroxide Mg3P2 Au(OH)3 Magnesium phosphide Gold (III) hydroxide Name nitric acid plumbic iodide Formula HNO3 PbI4 Neon gas iron (III) hydrogen sulfite Chlorine gas potassium bromite copper (II) sulfate nonahydrate manganese (II) hypochlorite hydrosulfuric acid Cl2 (g) KBrO2 CuSO4•9H2O BaO2 manganese dioxide sulfuric acid silver sulfide Barium peroxide MnO2 H2SO4 (aq) Ag2S lead (IV) perchlorate magnesium phosphite Ne (g) Fe(HSO3)2 Mn(ClO)2 H2S (aq) Pb(ClO4)4 Mg3(PO3)2 . P. NF3. 8. 7 13. I 20. same atomic # but different mass # 8. cathode 6. empty space. Ionic: electrostatic attraction between charged ions. C. Nap11n12 2. Covalent: equal sharing of a pair of electrons (no partial charges or dipoles). C 3. S.49 u 10. 43% X-59 11. lower nuclear charge) d) Cl (more nuclear charge) e) Cl (electrons closer to nucleus since fewer energy levels) f) Mg (weaker hold on electrons since more energy levels) 16. 8. 57% X-56.2. passed through. P. See Chemistry 11 textbook back cover 12. 1 9.SCH3U Exam Review 8 EXAM REVIEW ANSWERS: Parts 1 1. MM. a) 0.→ P3- 17. c) v Pp15n16 2. S 4. trigonal planar b) 0. BH3. X and Z. 1. unsat. 0.3 → 2. 288 g 15. 4 c) → calcium + sodium sulfate b) 2Ag + CuSO4 → Ag2SO4 + Cu (single displ) 23. 40. 6. limitng reactant questions) Make corrections to your unit tests and review multiple choice errors Study only 1 or 2 units at one sitting.02 mol b) 1. = disappears.g. weak acids do not.(aq) + 2K+ (aq) + 2I. K2CO3 4.77 mol e) 1. 30.1 b) LR = Cu2S c) 309 kg Cu d) 92.02 mol d) 1. no change = sat.30 b) 4. X > Z > W > Y EXAM REVIEW ANSWERS: Parts 2 1.8 g 7. Strong acids dissociate 100%. see notes 2. C6H6O2 5.9 mol/L 12. 1. 24. 2 L 36. 29..22 x 1024 molecules c) 2. 10 22. a) 2H2O → 2H2 + O2 (decomp) d) BaCl2 + Na2SO4 → BaSO4 + 2NaCl (double displ) c) CO2 + H2O → H2CO3 (syn) e) C4H10 + 15/2 O2 → 5CO2 + 5H2O (combust) f) NaOH + HNO3 → NaNO3 + H2O 24. 73 g/mol Studying Suggestions Create a word list for each unit Create a formula sheet for the course Create problem solving maps for major problems (e.6. take breaks while studying to avoid burning out Do the practice exam (available on-line at IRHS Science SC3U Æ Exam Review) Get lots of sleep and be sure to eat breakfast the morning of the exam AVOID excessive caffeine (e.(aq) → PbI2 (s) + 2K+ (aq) + 2NO3. 159 kpa 31. a) 2.28mol/L + 2+ + 216.67 L 14.SCH3U Exam Review 21.3 25. 23.44 L 27.3 e) 4. grows = supersat. 91. 0.22 Refer to notes. a) 10. 1 → 6.1% Na.3% O 3. 40. a)→ Na3PO4 (aq) + 3H2O b) →2NaI (aq) + CO2 (g) + H2O (l) 20. a) Pb(NO3)2 (aq) + 2KI (aq) → PbI2 (s) + 2KNO3 (aq)| b) Pb2+ (aq) + 2NO3.8 →5.3 c) 3. a) NR 9 b) 1. a) 84g/100g water b) 57oC c) 84 g precipitates out (all answers approximate) 11. 0.0 mol b) 1.20 mol/L 26. a) 2.02 mol f) 1. 0. a) → 2NaNO3 (aq) + Fe(OH)2 (s) 18. 29.2% yield 9.6% S.8 L 30.0124mol 33. coffee.0 L 8. 1. . a)KCl → K + Cl b) Ba(NO3)2 → Ba + 2NO3 c) (NH4)2SO3 →2NH4 + SO3 b) → 2AgCl(s) + Mg(NO3)2 (aq) 17.6 → 2.41 mol 34. 7 → 2.1. Red Bull) the night before or day of the exam AVOID pulling an “all nighter”! . Refer to notes. a) 1. 45.00 mol d) 937 g 6.4 b) → sodium fluoride + chlorine d) 1.g.10 → 1.1 kg 13.50 mol c) 3.(aq) c) Pb2+ (aq) + 2I. a) 2. 0.122 mol/L 28. 316 K 32. 10.(aq) → PbI2 (s) 19.1.757 g/L 35.1. a) 20% m/m b) 3.